Metallic Bonding and Properties of Structures

Giant Metallic Structures

1. The outermost shell of electrons of a metal atom is delocalised - the electrons are free to move about the metal. This leaves a positive metal ion, e.g. Na+, Mg2+, Al3+.

2. The positive metal ions are attracted to the delocalised negative electrons. They form a lattice of closely packed positive ions in a sea of delocalised electrons - this is metallic bonding.

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Metallic Properties

Metallic bonding explains why metals do what they do -

1. The number of delocalised electrons per atom affects the melting point. The more there are the stronger the bonding will be and the higher the melting point. Mg2+ has two delocalised electrons per atom, so it's got a higher melting point than Na+, which only has one. The size of the metal ion and the lattice structure also affect the melting point.

2. As there are no bonds holding specific ions together, the metal ions can slide over each other when the structure is pulled, so metals are malleable (can be shaped) and ductile (can be drawn into a wire).

3. The deolcalised electrons can pass kinetic energy to each other, making metals good thermal conductors.

4. Metals are good electrical condiuctors because the delocalised electrons can carry a current.

5. Metals are insoluble, except in liquid metals, because of the strength of the metallic bonds.

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Solids, Liquids and Gases Physical Properties

1. A typical solid has its particles close together. This gives it a high density and makes it incompressable. The particles vibrate about a fixed position and can't move about freely.

2. A typical liquid has a similar density to a solid and is virtually incompressible. The particles move about freely and randomly within the liquid, allowing it to flow.

3. In gases, the particles haveloads more energy and are much further apart. So the density is generally pretty low and it's very compressible. The particles move about freely, with not a lot of attraction between them, so they'll quickly diffuse to fill a container.

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Physical Properties of a Solid

Here are some handy points that'll make AS chemistry a little less painful:

1. Melting and boiling points depend on the attraction between particles.

2. The closer the particles, the greater the density.

3. If there are charged particles that free to move, then it'll conduct electricity.

4. Solubility depnds on they type of particles present.

Water is a polar solvent and it tends to only dissolve other polar substances.

5. If a solid has a regular structure, it's called a crystal. The structure is a crystal lattice.

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Covalent Bonds DON'T BREAK!!!!!

This is something that confuses loads of people - prepare to be enlightened.

1. To melt or boil a simple covalent compound you only have to overcome the Van der Waals or hydrogen bonds that hold the molecules together.

2. You don't need to break the much stronger covalent bonds that hold the atoms together in the molecules.

For example... Chlorine, Cl2, has a stronger covalent bonds than bromine, Br2. But under normal conditions, chlorine is a gas and bromine a liquid. Bromine has the highest boiling point because its molecules have more electrons, giving stronger Van der Waals forces.

Different Bonding Types

Observation matches Bonding Models

E.g. the physical properties of ionic compounds provide evidence that supports the theory of ionic bonding.

1. They have high melting points - this tells you that the atoms are held by a strong attraction. Positive and negative ions are strongly attracted, so the model fits the evidence.

2. They are often soluble in water but not in non-polar solvents - this tells you that the particles are charged. The ions are pulled apart by polar molecules like water, but not by non-polar molecules. Again, the model of ionic structures fits this evidence.

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Limitations

Like pretty much all models, bonding models aren't totally accurate.

1. Dot-and-cross models of ionic and covalent bonding are great for explaining what's happening nice and clearly. But like most things in life, it's not really quite as simple as that.

2. One important reason is that most bonds aren't purely ionic or purely covalent but somewhere in between. This is down to bond polarisation. Most compounds end up with a mixture of ionic and covalent properties.